Molecular Geometry. It’s all about the Electrons Electrons decide how many bonds an atom can have...
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Molecular Geometry
Molecular Geometry. It’s all about the Electrons Electrons decide how many bonds an atom can have They also decide the overall shape of the molecule OPPOSITES
Its all about the Electrons Electrons decide how many bonds an
atom can have They also decide the overall shape of the molecule
OPPOSITES ATTRACT!
Slide 3
Lewis Structures A Lewis structure is basically a diagram of
how a molecule looks using dots to represent the electrons. There
are 4 rules for making these structures and that is where the
electrons come into play.
Slide 4
Rule number 1 Count the Number of valence electrons! This means
of all the atoms present With a polyatomic anion, add one for each
negative charge With a polyatomic cation, subtract one for each
positive charge Ex: CO 2 C: 4 O: 6 6 + 6 + 4 = 16
Slide 5
Rule Number 2 Draw a skeleton structure for the molecule using
all single bonds This will most often be one central atom with
several surrounding ones Typically the central atom is written
first Ex: CO 2 O-C-O
Slide 6
Rule Number 3 Determine the number of valence electrons still
available for distribution To do this simply deduct two valence
electrons for each single bond written in step two Ex: CO 2 two
single bonds so far so we subtract a total of 4 16 4 = 12
Slide 7
Fourth Rule Determine the number of electrons required to fill
an octet for each atom If this equals the number of electrons left,
then place them on the atoms as unshared pairs If the number of
electrons available is less than the number needed then you need to
make double or triple bonds in place of the single bonds
Slide 8
CO 2 (again) O C O So far we have used four electrons so 16 4 =
12 Carbon still needs 4 more electrons and each Oxygen needs 6
more. 6 + 6 + 4 = 16 But we only have 12 left so lets make some
double bonds! O C O becomes O = C = O
Slide 9
Now Carbon doesnt need anymore electrons and the Oxygens only
need 4 more each. Since we used 4 electrons to make those into
double bonds we now have exactly 8 electrons left. Now we simply
distribute them to the Oxygen atoms as unshared paired
electrons.
Slide 10
Practice!
Slide 11
Resonance! Resonance is invoked whenever a single Lewis
structure does not adequately reflect the properties of a substance
In other words, resonance comes into play when you can make two
structures that are the same in their placement of atoms but
different in the bonds SO 2
Slide 12
Resonance structures are NOT forms where the electrons move
eternally between them Resonance structures are equally plausible
or they are not a resonance structure Resonance forms differ in
their distribution of electrons, NOT in their arrangement of atoms!
So just because a formula for a compound is the same it does not
mean that it is a resonance structure
Slide 13
Electronegativity Electronegativity is a measure of how much an
element wants to pull electrons towards itself This is represented
as a unit-less number ranging from 0 4.0 Heres a handy reference
sheet with all the values. Guard it with your LIFE!
Slide 14
So what? These numbers can be used mathematically to know if a
bond is ionic or covalent It can also tell you if a covalent bond
is more polar or less polar (more on polarity in a minute) So all
we have to do is subtract one from the other.
Slide 15
Example Fluorine has an electronegativity of 4.0 Sodium has an
electronegativity of 0.9 4.0 0.9 = 3.1 So what does that mean? It
means that it is an ionic bond! This makes sense since we know that
a bond involving one metal and one non- metal is ionic.
Slide 16
Example two Fluorine has an electronegativity of 4.0 Carbon has
an electronegativity of 2.15 4.0 2.5 = 1.5 This makes this bond
covalent!
Slide 17
Sharing is caring, but some elements are greedy! This
greediness shown by some elements like fluorine leads us to the
next piece of this puzzle The more unequal the sharing of electrons
is in a bond, the more polar it is. The smaller that difference in
electronegativity, the less polar.
Slide 18
Polar vs Non-polar Polar: Number greater than 0.4 Unequal
sharing of electrons Water is an example Non-Polar: Number less
than 0.4 Equal sharing of electrons Methane (CH 4 ) is an
example
Slide 19
So why is this important? Polarity is a major component of
organic chemistry Polarity also explains why certain substances can
dissolve other substances while others cannot Think oil and
water